Lewis Diagrams give us information about the arrangement and bond order of compounds.
Lewis Diagrams give us information about the arrangement and bond order of compounds (0:23). There are four steps to write Lewis Structure: 1) Sum up available valence electrons (1:16), 2) write symbols and connect them with a single bond (1:57), 3) complete the octet for surrounding atoms (2:27), and 4) add remaining electrons as lone pairs on central atom (2:48).
Some elements, like hydrogen, boron and beryllium as well as elements in Period 3 and beyond are exceptions to the octet rule (3:24). Formal charges can be used to determine the best Lewis Structure (5:15). In our episode we discuss calculating the formal charge for two possible structures of CO2 (6:07).
Which two elements will never be in the center of a Lewis Structure?
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap Lewis Diagrams and Formal Charges.
Lets Zoom out:
Unit 2 - Molecular and Ionic Compound Structure and Properties
Topic - 2.5 and 2.6 Lewis Structures and Formal Charge
Big idea - Structure and Properties
Introduction:
In episode 47, when talking about Empirical and Molecular formulas, we used formaldehyde, CH2O, as an example. So let’s stick with it and preserve some consistency. Chemical Formulas tell us either the ratio of atoms in the ionic compound or, as with formaldehyde, the actual number of covalently bonded atoms. But they give us very little information about the arrangement of atoms or the bond order. To determine these we need to draw the Lewis Diagrams, also known as Lewis Structures.
Lets Zoom in:
To write Lewis Structures for molecular compounds, we follow these steps: The first step is to sum the valence electrons available by adding up the valence electrons of each atom in the molecule. In our example, CH2O, we have 4 valence electrons from carbon, six valence electrons from oxygen and 2 times 1 electron from hydrogen, which adds up to 12 electrons. When counting the valence electrons, we also have to take into account the charges of the compound. For anions, we add an electron to the sum for each negative charge, for cations we subtract one electron for each positive charge.
In the second step, we write the symbols for the atoms. Generally, the more electropositive, or less electronegative, element goes into the center - the other atoms surround it. In our case, carbon is the central atom. We then connect all atoms to the center atom with a single bond, representing the two shared electrons. Keeping track of our electrons, we have used 3 times 2, therefore 6 electrons for these single bonds and still have 6 remaining.
The third step is to complete the octet for the atoms on the outside by adding them as lone pairs. This doesn’t apply to hydrogen, which will never have an octet. These lone pairs are represented as pairs of dots. In our example, we put the six remaining electrons on oxygen.Keeping count of our electrons, notice that we have no electrons remaining.
In a fourth step, we would add the remaining electrons as lone pairs to the central atom to complete its octet. But what if I don’t have enough electrons remaining to complete the octet, like in our case? That indicates that you have to use a double or triple bond, by pulling in one of the lone pairs of an atom to form a second or third bond between the surrounding atom and the central atom. In our example, we pull one lone pair from oxygen to be a double bond between carbon and oxygen. Now all atoms have the number of electrons needed.
Let’s briefly talk about the exceptions, because, as we all know, Chemistry is the Science of Exceptions: Aside from hydrogen, elements in the second period before carbon, that means Boron and Beryllium, can make stable compounds with fewer than eight electrons. These are electron-deficient compounds. Examples would be Boron trifluoride, BF3, and beryllium fluoride, BeF2.
Of course, we also have the other side of the spectrum: central atoms with more than eight valence electrons. These are the elements that are in period 3 or higher in the periodic table, for example phosphorus and sulfur. They can use their d orbitals to make more than four bonds. Therefore, you have compounds like phosphorus pentafluoride, PF5, and sulfur hexafluoride, SF6.
In some instances, more than one Lewis Structure is possible, like with CO2. Briefly pause this episode and take a moment to draw a possible structure for CO2 following our steps to draw Lewis Structures. OK, let’s take a look at the possible structures: You either have O double bond C double bond O or you have O single bond C triple bond O. The second one could, of course, also have the triple bond first and the single bond second. Both of these structures are valid: for both of them, the octet is fulfilled for all atoms. So, how do you know which Lewis Structure is better?
You can determine the formal charge for each atom, which is the charge the atom would have if all electrons forming the covalent bond were shared equally. It can be calculated for each atom of the compound using this formula: formula charge equals the number of valence electrons according to the periodic table minus half of the electrons in bonds minus all nonbonding electrons. Let’s look at the carbon dioxide with the double bonds. Oxygen, according to the Periodic Table, has 6 valence electrons. In the double-bonded carbon dioxide, it has 2 bonding electrons as well as four nonbonding electrons. 6 minus 2 minus 4 is zero. The carbon has 4 valence electrons minus 4 bonding electrons and therefore also a formal charge of 0.
It is different in the second possible structure. The oxygen with the single bond has one bonding electron and six non-bonding electrons. It’s formal charge is therefore six valence electrons minus 1 bonding electron minus 6 nonbonding electrons equals minus 1. On the other side, the oxygen with the triple bond has a formal charge of +1, since it has 6 valence electrons minus 3 bonding electrons, minus 2 nonbonding electrons. So which one is better? Here are the guidelines: the lewis structure with a formal charge closest to 0 is best. In our case that would be the Lewis Structure with two double bonds. What if both have formal charge? Then the better structure is the one where a negative formal charge is assigned to the more electronegative element.
In neutral compounds the formal charges cancel out and they have an overall formal charge of 0. Compounds with charges, for example polyatomic ions, have an overall formal charge that matches their ionic charge.
To recap…
Lewis Diagrams show us the arrangement and bond order of chemical compounds. They can be drawn in five steps: 1) Count available valence electrons. 2) Draw a skeleton structure by adding single bonds between central atoms and surrounding atoms. 3) Complete the octet for surrounding atoms. 4) Add remaining electrons to the central atom. Double and triple bonds might be necessary. Exceptions to the octet rule are for elements before carbon as well as beyond period 3. Calculating formal charges can help us to determine the best Lewis Structure.
Coming up next on the Apsolute RecAP Chemistry Edition: Unit 1 and 2: Selected FRQs.
Today’s Question of the day is about Lewis Structures:
Which two elements will never be in the center of a Lewis Structure?
C & N
H & F
P & S